AS Chemistry
Unit 3: Inorganic Chemistry
Section 1: Period 3
Part 1: Atomic and Physical Properties of the Elements
Atomic Properties
Task 1
Can you write the electronic configurations of the Period 3 elements in the
table below?
Na [Ne]
P
[Ne]
Mg [Ne]
S
[Ne]
Al
[Ne]
Cl
[Ne]
Si
[Ne]
Ar [Ne]
In each case, [Ne] represents the complete electronic structure of a neon atom.
Atomic radius
The graph shows how the atomic radius changes as you go across Period 3:
0.25
Atomic radius (nm)
0.2
0.15
0.1
0.05
0
Na
Mg
Al
Si
P
S
Cl
Ar
Element
Task 2
An atomic radius is a measure of the distance from the nucleus to the bonding
pair of electrons.
i)
What is the general trend in atomic radius across Period 3?
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Cambridge A-level Centre
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ii)
Unit 3: Inorganic Chemistry
Can you use your knowledge of atomic structure to explain this trend?......
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iii)
Why do you think the value for argon does not follow the expected
trend?....................................................................................................................................
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Ionic radius
The radii of the ions also decrease across a Period, but it must be remembered
that elements on the left form cations and elements on the right form anions.
ion
ionic
radius (nm)
+
Na
0.098
Cations
Mg
Al3+
0.065
0.045
2+
4+
Si
0.038
3-
P
0.212
Anions
S20.190
Cl0.181
Task 3
The radii of the cations are all smaller than that of the corresponding atoms,
but the radii of the anions are all larger. Can you explain why?
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Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
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First ionisation energy
First ionisation energy (kJ mol-1)
This has already been covered in Unit 1. You should be able to use your
scientific knowledge to explain in detail the trend shown in the graph below:
1600
1400
1200
1000
800
600
400
200
0
Na
Mg
Al
Si
P
S
Cl
Ar
Element
Remember, first ionisation energy is governed by:




the charge on the nucleus;
the distance of the outer electron from the nucleus;
the amount of screening by inner electrons;
whether the electron is alone in an orbital or one of a pair.
Electronegativity
Task 4
Can you write a definition for the term ‘electronegativity?
Electronegativity is .................................................................................................................
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Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
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The Pauling scale is the most commonly used. Fluorine (the most electronegative
element) is assigned a value of 4.0, and values range down to caesium and
francium which are the least electronegative at 0.7.
The trend
Pauling electronegativity value
The Electronegativity value for chlorine is given. Can you sketch the rest of the
graph?
3.5
3
2.5
2
1.5
1
0.5
0
Na
Mg
Al
Si
P
S
Cl
Ar
Element
Explaining the trend
The trend is explained in exactly the same way as the trend in atomic radii.
As you go across the period, the bonding electrons are always in the same level –
level 3. They are always being screened by the same inner electrons.
All that differs is the number of protons in the nucleus. As you go from sodium
to chlorine, the number of protons steadily increases and so attracts the
bonding pair more closely.
Task 5
Why is no value included for argon?....................................................................................
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
Physical Properties
Here we look at the electrical conductivity, and the melting and boiling points of the elements. To understand these, you
first have to understand the structure of each of the elements.
Task 6
Use your knowledge from Unit 1 to complete the table below.
Na
Structure
Type of element
Bonding
Formula
Type of force broken
on melting/boiling
Does the element
conduct electricity?
Cambridge A-level Centre
Mg
Al
Si
P
S
Cl
Ar
AS Chemistry
Unit 3: Inorganic Chemistry
Task 7
Use the information in the table above to answer the following questions.
1.
(a)
Explain why electrical conductivity decreases across Period 3
from sodium to phosphorus.
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(3)
(b) The table below shows the melting temperatures, Tm, of the
Period 3 elements.
Element
Na
Mg
Al
Si
P
S
Cl
Ar
Tm/K
371
923
933
1680
317
392
172
84
Explain the following in terms of structure and bonding.
(i)
Magnesium has a higher melting temperature than sodium.
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(ii)
(2)
Silicon has a very high melting temperature.
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(2)
(iii) Sulphur has a lower melting temperature than magnesium.
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(iv) Argon has a lower melting temperature than chlorine.
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(2)
AS Chemistry
Unit 3: Inorganic Chemistry
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(3)
(Total 14 marks)
References
A-level Chemistry pages 197-207
Chemistry in Context pages 42-47, 169-176
Learning Objectives
Candidates should be able to:
 describe qualitatively (and indicate the periodicity in) the variations in
atomic radius, ionic radius, melting point and electrical conductivity of the
elements (see the Data Booklet)
 explain qualitatively the variation in atomic radius and ionic radius
 interpret the variation in melting point and in electrical conductivity in
terms of the presence of simple molecular, giant molecular or metallic
bonding in the elements
 explain the variation in first ionisation energy.
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
Section 1: Period 3
Part 2: Reactions with oxygen
Reactions of Period 3 Elements with oxygen
As we move across a period of the Periodic Table from left to right, we can see
that there are small, but regular changes in atomic structure. These small
changes can produce very big variations in the chemical and physical properties
of the elements and their compounds. The pattern is then repeated as you go
across the next period. The occurrence of periodic patterns is called
Periodicity. The physical properties of the Period 3 elements have been
discussed previously. Here we look at their chemical properties.
Task 1
Can you complete the table below?
Group number
1
2
3
4
5
6
Element in Period 3
Nuclear charge
[Ne] electronic
configuration
Trend in Atomic radius
Trend in 1st ionisation
energy
Trend in
electronegativity
Formula of oxide/s
The reactivity and properties of elements depend upon a combination of things:
nuclear charge, size of the atom, the number of outer electrons and the amount
of shielding, and these will also help to explain and predict the properties of
many compounds. It is therefore very useful if you are familiar with the trends
shown in the table above. You should not, however, lose sight of your general
chemistry knowledge.
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Unit 3: Inorganic Chemistry
Task 2
Can you complete the notes below by adding the missing words and writing the
appropriate balanced equations?
With oxygen
The solid elements in Period 3 all burn in air or oxygen when ignited. Sodium
burns with a __________ flame, forming the oxide:
Magnesium, aluminium, silicon and phosphorus burn when ignited, emitting a very
bright white light and white smoke of the oxides:
These reactions are all very exothermic.
Sulphur burns with a __________ flame but much less vigorously than the
elements above, to form the pungent, colourless gas ____________________:
In an excess of pure oxygen, some SO3 is also formed. This utilises the highest
oxidation state of sulphur.
Task 3
Can you complete the table below?
Tm/K
Bonding
Structure
Na2O
1548
MgO
3125
Al2O3
2345
SiO2
1883
P4O10
573
Use your scientific knowledge to explain the changes in melting point.
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SO2
200
AS Chemistry
Unit 3: Inorganic Chemistry
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Acid-base properties of the oxides of Period 3 Elements
The table below shows the change in pH for solutions of the oxides across
Period 3. There is an evident trend for alkaline oxides
→ acidic oxides as
the bonding changes from ionic to covalent, but it is masked by the change in
solubilities. A substance will only change the pH of water if it dissolves.
Task 4
Can you write balanced equations for the reactions of each oxide with water to
give the appropriate pH value.
Oxide
Reaction with water
Na2O
pH
14
MgO
9
Al2O3
7
SiO2
7
P4O10
0
SO2
3
SO3
0
Ionic oxides
The oxide ion is too highly charged to exist on its own in water. It attracts
water molecules and hydrolyses to form OH-(aq) ions:
O2− + H2O → 2 OH−
Cambridge A-level Centre
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Unit 3: Inorganic Chemistry
Na2O: very soluble in water → lots of oxide ions present and therefore many
OH-(aq) ions will be formed.
MgO: less soluble in water due (in part) to higher lattice energy → less oxide
ions and less OH-(aq) ions.
Al2O3: insoluble → no oxide ions will be present and hence no OH-(aq) ions will be
formed. Aluminium oxide is an ionic solid with some covalent character. It is
amphoteric; it will act as both an acid and a base.
BASE: Al2O3
+
3H2SO4
→
Al2(SO4)3
+
ACID: Al2O3
+
2NaOH
+
3H2O
2NaAl(OH)4
→
3H2O
Covalent oxides
SiO2: insoluble in water → no ions produced.
P4O10, SO3:
these acidic oxides rapidly dissolve in water and give the
corresponding acids. They are very soluble, so many H+(aq) ions are formed, giving
a low pH.
SO2 is less soluble so a weaker acid is formed.
References
A-level Chemistry pages 209-211
Chemistry in Context pages 177-184
Learning Objectives
Candidates should be able to:
 describe the reactions, if any, of the elements with oxygen to give Na2O,
MgO, Al2O3, P4O10, SO2 and SO3.
 state and explain the variation in oxidation number of the oxides.
 describe the reactions of the oxides with water.
 describe and explain the acid/base behaviour of oxides and hydroxides,
including, where relevant, amphoteric behaviour in reaction with NaOH
and acids.
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Unit 3: Inorganic Chemistry
Section 1: Period 3
Part 3: Reactions with chlorine
All of the chlorides of the period 3 elements can be formed by direct combination with chlorine but your syllabus is only
interested in the reactions of Na, Mg, Al, Si and P.
Task 1
Use the information on pages 47-49 and 179-180 of ‘Chemistry in Context’ and pages 209 and 211-212 of ‘AS level
Chemistry’ to complete the table below:
Element
Description of reaction
with chlorine
Formula of chloride/s
Oxidation state of
period 3 element
State of chloride at
r.t.p.
b.pt. of chloride (oC)
Structure of chloride
Bonding in chloride
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Na
Mg
Al
Si
P
AS Chemistry
Unit 3: Inorganic Chemistry
Task 2
Read the information sheet ‘Aluminium chloride’ and answer the question below.
i. At room temperature aluminium chloride is considered to be ionic. What two
pieces of evidence seem to contradict this?
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ii. At temperatures above ~200oC aluminium chloride exists as a dimer, Al2Cl6.
This molecule contains both ordinary and dative covalent bonds. Draw a
diagram to show the two types of bonding.
Reaction of the period 3 chlorides with water
As Period 3 is crossed, the reactions of the chlorides with water become
increasingly more violent.


Ionic chlorides usually dissolve in water to form neutral solutions
containing the hydrated parent ions.
Covalent chlorides are hydrolysed by water to form acidic solutions
containing HCl.
Cambridge A-level Centre
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Unit 3: Inorganic Chemistry
Task 3
Can you summarise these reactions in the table below?
Reaction with water
pH
7
6/7
3
0
0
References
A-level Chemistry pages 209-210 and 211- 212
Chemistry in Context pages 47-49 and 179-180
Learning Objectives
Candidates should be able to:
 describe the reactions, if any, of the elements with chlorine to give NaCl,
MgCl2, Al2Cl6, SiCl4, and PCl5.
 describe and explain the reactions of the chlorides with water.
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Section 2: Group 7
Unit 3: Inorganic Chemistry
Part 1: Properties of the elements
Please note that fluorine will not be included in discussions of Group VII,
because its small size and high electronegativity give it some anomalous
properties.
Task 1
Use your general scientific knowledge to fill in the gaps in the paragraph below.
The halogens are the __________ in _______7 of the __________ Table. All
halogen atoms have 7 __________ in the outer shell. The halogens are the
most __________ group of non-metals, and none of them is found
__________ in the elemental form. They are all found in __________, often
as __________ ions (a singly __________ charged ion, e.g. Br-).
Fluorine and chlorine are the most __________ halogens, bromine occurs in
smaller __________, iodine is quite __________ and astatine is an
__________ produced, short-lived, __________ element.
All the halogen elements occur as __________ molecules, e.g. F2. The atoms
are linked by a single __________ bond.
In compounds, a halogen atom can attain stability by:

__________ an electron from a __________ atom to form a halide ion
in an __________ bonded compound,

sharing an electron from another atom in a __________ bonded
compound.
Task 2
Draw electron dot-cross diagrams for NaCl and HCl in the space below.
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
Task 3
Use the information in your textbooks to complete the table below.
Element
Electronic
structure
Appearance
at r.t.p.
State
r.t.p.
at
Melting
point (oC)
Boiling point
(oC)
Colour
water
in
Colour
hexane
in
Fluorine
Chlorine
Bromine
Iodine
Task 4
From your knowledge of the structure and bonding of the halogens, explain why
they are more soluble in organic solvents than in aqueous solution.
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Trends in Melting Point and Boiling Point
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
You will see that both melting points and boiling points rise as you descend the
Group.
Task 5
Use your scientific knowledge to explain the trend in volatility of the halogens.
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Trends in chemical properties
Task 6
Watch the short video extract. For each reaction write down your observations.
Halogen
Chlorine, Cl2
Halide ion in aqueous solution
Chloride, ClBromide, Br-
Iodide, I-
Bromine, Br2
Iodine, I2
The oxidising power of the halogens
Imagine the reaction between one halogen (chlorine, say) and the ions of
another one (iodide ions, perhaps). The iodide ions will be in a solution of a salt
like sodium or potassium iodide. The sodium or potassium ions will be spectator
ions, and are completely irrelevant to the reaction. In the chlorine and iodide ion
case, the reaction would be:
Cl2 + 2I-  2Cl- + I2
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
The iodide ions have lost electrons to form iodine molecules. They have been
__________.
The chlorine molecules have gained electrons to form chloride ions. They have
been __________.
This is obviously a redox
_______________ agent.
reaction
in
which
chlorine
is
acting
as
Fluorine
We'll have to exclude fluorine from this descriptive bit, because it is too strong
an oxidising agent. Fluorine oxidises water to oxygen and so it is impossible to
do simple solution reactions with it.
Chlorine, bromine and iodine
In each case, a halogen higher in the Group can oxidise the ions of one lower
down. For example, chlorine can oxidise the bromide ions (in, for example,
potassium bromide solution) to bromine:
The bromine appears as an orange solution.
As you have seen above, chlorine can also oxidise iodide ions (in, for example,
potassium iodide solution) to iodine:
The iodine appears either as a red/brown solution if you are mean with the
amount of chlorine you use, or as a dark grey precipitate if the chlorine is in
excess.
Bromine can only oxidise iodide ions to iodine. It isn't a strong enough oxidising
agent to convert chloride ions into chlorine. (You have just seen exactly the
reverse of that happening.)
A red/brown solution of iodine is formed (see the note above) until the bromine
is in excess. Then you get a dark grey precipitate.
Cambridge A-level Centre
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Unit 3: Inorganic Chemistry
Iodine won't oxidise any of the other halide ions (unless you happened to have
some extremely radioactive and amazingly rare astatide ions - astatine is at the
bottom of this Group).
To summarise




Oxidation is loss of electrons. Each of the elements (for example,
chlorine) could potentially take electrons from something else to make
their ions (e.g. Cl-). That means that they are all potentially oxidising
agents.
Fluorine is such a powerful oxidising agent that you can't reasonably do
solution reactions with it.
Chlorine has the ability to take electrons from both bromide ions and
iodide ions. Bromine and iodine can't get those electrons back from the
chloride ions formed. That means that chlorine is a more powerful
oxidising agent than either bromine or iodine.
Similarly bromine is a more powerful oxidising agent than iodine. Bromine
can remove electrons from iodide ions to give iodine - and the iodine can't
get them back from the bromide ions formed.
This all means that oxidising ability falls as you go down the Group.
Relative
oxidising
power
Fluorine
Chlorine
Bromine
Iodine
Reactions of the halogens with hydrogen
The halogens all react with hydrogen to give the corresponding hydrogen halide.
The relative reactivity of the halogens is well illustrated by these reactions:




Fluorine explodes with hydrogen even in the dark at -200oC.
The reaction with chlorine is explosive when exposed to ultraviolet light.
With bromine, the reaction occurs slowly on heating.
The reaction with iodine is incomplete on heating.
Cambridge A-level Centre
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Unit 3: Inorganic Chemistry
Task 7
Write equations, including state symbols, for the reactions of chlorine, bromine
and iodine with hydrogen gas.
The trend is further illustrated by consideration of the ΔHfθ values:
ΔHfθ
-1
(kJ mol )
HCl
-92
HBr
-36
HI
+26
The hydrogen halides formed are all simple molecular compounds which are
gaseous at room temperature. As the size of the halogen atom increases down
the Group, the H-X bond length also increases. Consequently, the H-X bond
enthalpy decreases down the Group and the hydrogen halides become less
stable:



hydrogen chloride is stable at 1500oC
hydrogen bromide decomposes appreciably at 800oC, and
hydrogen iodide decomposes appreciably at 500oC.
If a red hot glass rod is placed in a gas jar containing hydrogen iodide, purple
fumes of iodine are seen. This shows that iodide ions are easily oxidised and can
act as strong reducing agents.
References
A-level Chemistry pages 237-240
Chemistry in Context pages 227-231
Learning Objectives
Candidates should be able to:
 describe the trends in volatility and colour of chlorine, bromine and
iodine.
 interpret the volatility of the elements in terms of van der Waals’ forces.
 describe the relative reactivity of the elements as oxidising agents.
 describe and explain the reactions of the elements with hydrogen.
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AS Chemistry


Unit 3: Inorganic Chemistry
describe and explain the relative thermal stabilities of the hydrides.
interpret these relative stabilities in terms of bond energies.
Cambridge A-level Centre
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Section 2: Group 7
Unit 3: Inorganic Chemistry
Part 2: Properties of the halides
Here we look at the redox reactions involving halide ions and concentrated
sulphuric acid. We use these reactions to discuss the trend in reducing ability
of the ions as you go from fluoride to chloride to bromide to iodide.
The Facts
There are two different types of reaction which might go on when concentrated
sulphuric acid is added to a solid ionic halide like sodium fluoride, chloride,
bromide or iodide. The concentrated sulphuric acid can act both as an acid and
as an oxidising agent.
Concentrated sulphuric acid acting as an acid
The concentrated sulphuric acid gives a hydrogen ion to the halide ion to
produce a hydrogen halide. Because this is a gas, it immediately escapes from
the system. If the hydrogen halide is exposed to moist air, you see it as steamy
fumes.
As an example, concentrated sulphuric acid reacts with solid sodium chloride in
the cold to produce hydrogen chloride and sodium hydrogensulphate.
All of the halide ions (fluoride, chloride, bromide and iodide) behave similarly.
The reduction of concentrated sulphuric acid
With fluoride or chloride ions
The fluoride and chloride ions aren't strong enough reducing agents to reduce
the sulphuric acid.
Whichever way you look at it, all you get is the hydrogen halide!
That isn't true, though, with bromides and iodides.
With bromide ions
The bromide ions are strong enough reducing agents to reduce the concentrated
sulphuric acid to sulphur dioxide gas. In the process the bromide ions are
oxidised to bromine.
Cambridge A-level Centre
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Unit 3: Inorganic Chemistry
Task 1
Can you write two electron half equations for these steps, then combine them to
form the overall redox equation for the reaction.
Task 2
Complete the notes below.
The bromide ions reduce the sulphuric acid to sulphur dioxide gas. This is a
decrease of oxidation state of the sulphur from _____ in the sulphuric acid to
_____ in the sulphur dioxide.
What you see in this reaction are the steamy fumes of __________
__________ contaminated with the brown colour of __________ vapour. The
sulphur dioxide is a colourless gas, so you couldn't observe its presence directly.
With iodide ions
Iodide ions are stronger reducing agents than bromide ions are. They are
oxidised to iodine by the concentrated sulphuric acid. The reduction of the
sulphuric acid is more complicated than before. The iodide ions are powerful
enough reducing agents to reduce it



first to sulphur dioxide
then to sulphur itself
and all the way to hydrogen sulphide.
Task 3
Can you give the oxidation state of sulphur in each of the species above?
The most important of this mixture of reduction products is probably the
hydrogen sulphide.
Cambridge A-level Centre
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Unit 3: Inorganic Chemistry
Task 4
Can you write the half equations for these reactions and combine them to give
the overall redox equation?
This time what you see is a trace of steamy fumes of hydrogen iodide, but
mainly lots of iodine. The reaction is exothermic and so purple iodine vapour is
formed, and probably dark grey solid iodine condensing around the top of the
tube.
You won't see the colourless hydrogen sulphide gas, but might pick up its "bad
egg" smell if you were foolish enough to smell the intensely poisonous gases
evolved!
Summary of the trend in reducing ability




Fluoride and chloride ions won't reduce concentrated sulphuric acid.
Bromide ions reduce the sulphuric acid to sulphur dioxide. In the process,
the bromide ions are oxidised to bromine.
Iodide ions reduce the sulphuric acid to a mixture of products including
hydrogen sulphide. The iodide ions are oxidised to iodine.
Reducing ability of the halide ions increases as you go down the Group.
(Obviously, this is the opposite direction to their oxidising ability!!)
Fluoride
Relative
reducing
power
Chloride
Bromide
Iodide
Cambridge A-level Centre
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Unit 3: Inorganic Chemistry
Testing for halide ions
Using silver nitrate solution
Carrying out the test
This test has to be done in solution. If you start from a solid, it must first be
dissolved in pure water.
The solution is acidified by adding dilute nitric acid. (Remember: silver nitrate +
dilute nitric acid.) The nitric acid reacts with, and removes, other ions that
might also give a confusing precipitate with silver nitrate.
Silver nitrate solution is then added.
Task 5
Can you complete the table below to show what you would observe in each case?
ion present
observation
FClBrIAll of the precipitates change colour if they are exposed to light - taking on
grey or purplish tints.
The absence of a precipitate with fluoride ions doesn't prove anything unless
you already know that you must have a halogen present and are simply trying to
find out which one. All the absence of a precipitate shows is that you haven't
got chloride, bromide or iodide ions present.
The chemistry of the test
The precipitates are the insoluble silver halides - silver chloride, silver bromide
or silver iodide.
Cambridge A-level Centre
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Unit 3: Inorganic Chemistry
Silver fluoride is soluble, and so you don't get a precipitate.
Confirming the precipitate using ammonia solution
Carrying out the confirmation
Ammonia solution is added to the precipitates.
Task 6
Can you complete the table below to show what you would observe in each case?
original precipitate
observation
AgCl
AgBr
AgI
The chemistry of the test
When the precipitate of silver chloride or silver bromide dissolves, the silver
ion forms a complex ion:
E.g.
AgCl(s)
+
2NH3(aq)
→
[Ag(NH3)2]+(aq)
+
Cl-(aq)
The reactions of chlorine with sodium hydroxide
Chlorine is used in the manufacture of bleach, which is widely used as a
disinfectant. Bleach is produced by passing chlorine gas up a tower, down which
cold, dilute aqueous sodium hydroxide is flowing. The equation for the reaction
which takes place in the tower is as follows:
Cl2(g) +
2NaOH(aq)
→
NaClO(aq)
+
Sodium chlorate (I)
Cambridge A-level Centre
NaCl(aq)
+ H2O(l)
AS Chemistry
Unit 3: Inorganic Chemistry
Task 7
i. Using oxidation states can you identify which element is oxidised and which
is reduced in the equation above?
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ii. What word do we use to describe reactions of this sort?
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With hot, concentrated aqueous sodium hydroxide a similar reaction occurs but
this time a mixture of chloride and chlorate (V) ions is formed.
Task 8
Can you write a balanced equation for the reaction above?
The sodium chlorate (V) produced is used as a weedkiller.
Task 9
Use pages 242-243 of your AS level textbook to list the uses of the halogens
below. Add any chemical formulae or equations you think may be useful.
Uses of the
halogens
Cambridge A-level Centre
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Unit 3: Inorganic Chemistry
References
A-level Chemistry pages 240-242
Chemistry in Context pages 233-236
Learning Objectives
Candidates should be able to:
 describe and explain the reactions of halide ions with
o aqueous silver ions followed by aqueous ammonia
o concentrated sulphuric acid.
 describe and interpret in terms of changes of oxidation number
the reaction of chlorine with cold, and with hot, aqueous sodium
hydroxide.
 explain the use of chlorine in water purification.
 recognise the industrial importance and environmental
significance of the halogens and their compounds, (e.g. for
bleaches; PVC; halogenated hydrocarbons as solvents,
refrigerants and in aerosols).
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
Section 3: Nitrogen and Sulphur
Nitrogen
Use pages 259-264 of ‘AS level Chemistry’ to find the answers to the following
questions?
1. Nitrogen gas makes up what percentage of the Earth’s atmosphere?
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2. Draw a dot-cross diagram to show the bonding in a molecule of nitrogen:
3. Nitrogen is used as an inert atmosphere in food storage and during some
metalworking processes where oxidation must be prevented. Why do you
think nitrogen gas is so unreactive?
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4. Can you list two situations in which nitrogen gas will react with oxygen gas?
Why do these reactions occur?
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Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
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5. Atmospheric nitrogen is ‘fixed’ as nitrate ions by bacteria in the roots of
certain plants such as peas, beans and clover. How can bacteria enable this to
happen?
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6. Nitrogen is converted into ammonia in the Haber Process. Can you write a
balanced equation for this reaction.
7. Ammonia readily forms ammonium ions through co-ordinate bonds. Can you
draw a diagram to represent this below:
Ammonia is acting as what type of species in this reaction?
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8. How would you test for the presence of an ammonium ion?
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Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
9. Much ammonia is converted into salts for use as fertilisers. What is a
fertiliser?
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Can you name three of these salts?
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10. Ammonia is also converted into nitric acid. Can you write a balanced equation
for the overall reaction and give the conditions needed?
11. There are concerns over the use of nitrate fertilisers as they have been
linked to the process of eutrophication. Can you draw a flow diagram for this
process below:
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
Sulphur
Use your scientific knowledge and the information on pages 262-264 of your
textbook to help you answer the following questions.
1. The Earth’s atmosphere contains oxides of nitrogen and sulphur dioxide from
natural sources. Can you name some of these natural sources?
a. nitrogen oxides........................................................................................................
b. sulphur dioxide........................................................................................................
2. Sulphur dioxide in the atmosphere has been linked to human respiratory
problems such as bronchitis and asthma. It is however, used as a food
preservative. How does sulphur dioxide preserve food?
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3. Sulphur dioxide is oxidised in the atmosphere by nitrogen dioxide. Write a
balanced equation for this reaction.
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4. Acid rain is corrosive and will attack buildings made of limestone and marble.
Why are these buildings so vulnerable?
5. What is being done to reduce sulphur dioxide emissions?
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6. Sulphur extracted from crude oil is used to make sulphuric acid in the
Contact Process. Can you give four uses for sulphuric acid?
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Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
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References
A-level Chemistry pages 259-264
Learning Objectives
Candidates should be able to:
 explain the lack of reactivity of nitrogen.
 describe the
o formation, and structure of, the ammonium ion
o the displacement of ammonia from its salts.
 understand the environmental consequences of the uncontrolled use of
nitrate fertilisers.
 understand and explain the occurrence, and catalytic removal, of oxides
of nitrogen.
 explain why atmospheric oxides of nitrogen are pollutants, including their
catalytic role in the oxidation of atmospheric sulphur dioxide.
 describe the formation of atmospheric sulphur dioxide from the
combustion of sulphur contaminated carbonaceous fuels.
 state the role of sulphur dioxide in the formation of acid-rain and
describe the main environmental consequences of acid-rain.
 understand the industrial importance of sulphuric acid.
 Describe the use of sulphur dioxide in food preservation.
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
Section 4: Group 2
Section 1: Atomic and physical properties
Atomic Radius
You can see that the atomic radius increases as you go down the Group.
Explaining the increase in atomic radius
Task 1
Which factors affect the size of the atomic radius?
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Compare beryllium and magnesium:
Be
1s22s2
Mg
1s22s22p63s2
In each case, the two outer electrons feel a core charge of 2+ from the nucleus.
The positive charge on the nucleus is shielded by the inner electrons.
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
This is equally true for all the other atoms in Group 2. Work it out for calcium if
you aren't convinced.
The only factor which is going to affect the size of the atom is therefore the
number of layers of inner electrons which have to be fitted in around the atom.
Obviously, the more layers of electrons you have, the more space they will take
up - electrons repel each other. That means that the atoms are bound to get
bigger as you go down the Group.
Trends in First Ionisation Energy
First ionisation energy is the energy needed to remove the most loosely held
electron from each of one mole of gaseous atoms to make one mole of singly
charged gaseous ions - in other words, for 1 mole of this process:
X(g)

X+(g)
+
e-
Notice that first ionisation energy falls as you go down the group.
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
Task 2
Can you explain this trend?
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Trends in Electronegativity
Electronegativity is a measure of the tendency of an atom to attract a bonding
pair of electrons. All of these elements have a low electronegativity (due to
their small core charge).
As the metal atoms get bigger, any bonding pair gets further and further away
from the metal nucleus, and so is less strongly attracted towards it. In other
words, as you go down the Group, the elements become less electronegative.
Cambridge A-level Centre
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Unit 3: Inorganic Chemistry
As you go down the Group, the bonds formed between these elements and other
things such as chlorine become more and more ionic. The bonding pair is
increasingly attracted away from the Group 2 element towards the chlorine (or
whatever).
References
A-level Chemistry pages 215-218
Chemistry in Context pages 211-213
Learning Objectives
Candidates should be able to interpret, and make predictions from, the trends
in physical and chemical properties of the elements and their compounds.
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AS Chemistry
Section 4: Group 2
Unit 3: Inorganic Chemistry
Section 2: Chemical properties
When Group 2 metals react, they are oxidised from a state of 0 to +2, forming
M2+ ions. The elements are powerful reducing agents. Group 2 metals become
increasingly reactive as you go down the Group. This is because the distance
between the nucleus and the outer electrons increases and so they become
easier to remove. Beryllium is markedly different from the other members of
the group and will not be considered in detail.
Reactions of the metals with water
Beryllium
Beryllium has no reaction with water or steam even at red heat.
Magnesium
Magnesium burns in steam to produce magnesium oxide and hydrogen.
Very clean magnesium has a very slight reaction with cold water. The reaction
soon stops because the magnesium hydroxide formed is almost insoluble in water
and forms a barrier on the magnesium preventing further reaction.
Task 1
Can you write equations for the reaction of Mg with both cold and hot water?
As a general rule, if a metal reacts with cold water, you get the metal
hydroxide. If it reacts with steam, the metal oxide is formed. This is because
the metal hydroxides thermally decompose (split up on heating) to give the
oxide and water.
Calcium, strontium and barium
These all react with cold water with increasing vigour to give the metal
hydroxide and hydrogen. Strontium and barium have reactivities similar to
lithium in Group 1 of the Periodic Table.
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
The equation for the reactions of any of these metals would be:
The hydroxides aren't very soluble, but they get more soluble as you go down
the Group. The calcium hydroxide formed shows up mainly as a white precipitate
(although some does dissolve). You get less precipitate as you go down the Group
because more of the hydroxide dissolves in the water.
Reactions of the metals with oxygen
The group 2 elements all burn in oxygen to form white solid oxides.
2M(s) +
O2(g) →
2MO(s)
Beryllium is reluctant to burn unless it is in the form of dust or powder. Its
flame colour is white. Beryllium has a very strong (but very thin) layer of
beryllium oxide on its surface, and this prevents any new oxygen getting at the
underlying beryllium to react with it.
Magnesium burns with a brilliant white flame. The others burn with
characteristic flame colours:
Calcium:
brick red
Strontium: crimson red
Barium:
apple green
Reactions of the oxides with water
The metal oxides are all basic and each reacts with water, forming a solution of
the hydroxide which is alkaline. The reaction of magnesium oxide with water is
slow; the other oxides react readily with water.
The solubilities of the hydroxides formed increase down the Group and the
solutions become more alkaline (approx. pH9 – 12).
Thermal stability of Group 2 carbonates and nitrates
Thermal decomposition is when a substance breaks down (decomposes) when
heated. The more thermally stable a substance is, the more heat it will take to
break it down.
You are only expected to describe the trend in thermal stability, not explain it.
Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
The thermal stability increases down the group.
Group 2 carbonates decompose to from the oxide and carbon dioxide. Group 2
nitrates decompose to form the oxide, nitrogen dioxide and oxygen:
Task 2
Can you write equations for the thermal decomposition of calcium carbonate and
calcium nitrate?
Uses
Task 3
Compounds of the Group 2 elements have many uses. Use the information spread
throughout your textbooks to answer the following questions.
1. Magnesium is the most commonly used metal in Group 2. Can you list 2
uses below?
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2. Which compounds of magnesium are used for the following? In each case
explain why.
a. Toothpaste and indigestion remedies......................................................................
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b. Furnace linings...............................................................................................................
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Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
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3. Which 3 types of rock contain large amounts of calcium carbonate?
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4. Several calcium compounds have common names based on the word ‘lime’.
Can you draw a flow diagram in the space below to show the relationship
between limestone, lime (or quicklime), slaked lime and limewater. Include
the chemical name and formula for each substance.
5. Find as many uses as you can for each of the following compounds of
calcium:
a. calcium carbonate.........................................................................................................
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b. calcium oxide.................................................................................................................
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Cambridge A-level Centre
AS Chemistry
Unit 3: Inorganic Chemistry
c. solid calcium hydroxide...............................................................................................
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d. calcium hydroxide solution.........................................................................................
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References
A-level Chemistry pages 217-222
Chemistry in Context pages 213-219
Learning Objectives
Candidates should be able to
 describe the reactions of the elements with oxygen and water.
 describe the behaviour of the oxides with water.
 describe the thermal decomposition of the nitrates and carbonates.
 interpret, and make predictions from, the trends in chemical properties
of the elements and their compounds.
 explain the use of magnesium oxide as a refractory lining material and
calcium carbonate as a building material.
 describe the use of lime in agriculture.
Cambridge A-level Centre